Question

Assign the oxidation state for nitrogen in each of the following. a. ${\mathrm{Li}}_3\mathrm{N}$ b. ${\mathrm{NH}}_3$ $\mathbf{c}.{\mathrm{N}}_2{\mathrm{H}}_4$ d. NO e. ${\mathrm{N}}_2\mathrm{O}$ $\mathbf{f}.{\mathrm{NO}}_2$ g. ${\mathrm{NO}}_2^{-}$ h. ${\mathrm{NO}}_3^{-}$ $\mathbf{i}.{\mathbf{N}}_2$

Short answer

The oxidation states of nitrogen in the given compounds are: a. Li3N: -3 b. NH3: -3 c. N2H4: -2 d. NO: +2 e. N2O: +1 f. NO2: +4 g. NO2-: +3 h. NO3-: +5 i. N2: 0

Step-by-step solution

Understand the rules for assigning oxidation states

To assign oxidation states to elements in compounds, we must follow these rules: 1. The oxidation state of an element in its elementary form (e.g., N2) is 0. 2. The oxidation state of an ion is equal to its charge. For example, in Li+ the charge is +1, so the oxidation state is +1. 3. The sum of oxidation states of all elements in a compound must be equal to the charge of the compound, which is 0 for neutral compounds and equal to the net charge for ions. 4. In some compounds, elements like hydrogen (H) and oxygen (O) have usual oxidation states: +1 for H and -2 for O, except in the cases of hydrides (H with oxidation state -1) and peroxides (O with oxidation state -1).

Assign oxidation states for each compound

a. Li3N In lithium nitride, Li has an oxidation state of +1 (Rule 2) and there are three Li ions, so the total positive charge is +3. To balance the charge (Rule 3), the oxidation state of nitrogen (N) must be -3. b. NH3. In ammonia, H has an oxidation state of +1 (Rule 4). Since there are three H atoms, the total positive charge is +3. Therefore, the oxidation state of nitrogen (N) must be -3 to balance the charge. c. N2H4. In hydrazine, H has an oxidation state of +1 (Rule 4). Since there are four H atoms, the total positive charge is +4. Because the compound is neutral (Rule 3) and there are two N atoms, each N atom must have an oxidation state of -2. d. NO. In nitrogen monoxide, oxygen (O) has an oxidation state of -2 (Rule 4). To balance the charge (Rule 3), the oxidation state of nitrogen (N) must be +2. e. N2O. In dinitrogen monoxide, oxygen (O) has an oxidation state of -2 (Rule 4). As the compound is neutral, the sum of oxidation states of the two N atoms must be +2 (Rule 3). Therefore, each N atom has an oxidation state of +1. f. NO2. In nitrogen dioxide, oxygen (O) has an oxidation state of -2. Since there are two O atoms, the total negative charge is -4. To balance the charge (Rule 3), the oxidation state of nitrogen (N) must be +4. g. NO2-. In nitrite ion (NO2-), oxygen (O) has an oxidation state of -2, and there are two O atoms, for a total negative charge is -4. Since the overall charge of the ion is -1 (Rule 3), the oxidation state of nitrogen (N) must be +3. h. NO3-. In nitrate ion (NO3-), oxygen (O) has an oxidation state of -2, and there are three O atoms, so the total negative charge is -6. Since the overall charge of the ion is -1 (Rule 3), the oxidation state of nitrogen (N) must be +5. i. N2. In diatomic nitrogen, both N atoms are in their elementary form (Rule 1), so the oxidation state for each nitrogen atom is 0.

Key concepts

Assigning Oxidation Numbers

Understanding how to assign oxidation numbers to elements within chemical compounds is vital for students studying chemistry. Oxidation numbers, also known as oxidation states, denote the degree of oxidation of an atom in a chemical compound.

Oxidation numbers are not necessarily the same as the actual charge on an atom, but they are useful for keeping track of electrons during chemical reactions, especially redox reactions. When assigning oxidation numbers, the following rules should be remembered:

• The oxidation state of a pure element is always zero.
• For monatomic ions, the oxidation state is equal to the ion's charge.
• In a molecule or compound, the sum of the oxidation states must equal the overall charge.
• Hydrogen generally has an oxidation state of +1, while oxygen typically has an oxidation state of -2. However, there are exceptions like in hydrides and peroxides.
• In compounds, other trends to consider include the typical oxidation states of elements based on their position in the periodic table.

Incorporating these rules will aid students in systematically determining each element's oxidation state in complex compounds.

Chemical Compounds Oxidation

Chemical compounds undergo oxidation when they lose electrons during a reaction, increasing the oxidation state of the chemical species involved. When we look at oxidation in compounds, we're often dealing with redox reactions, which involve the transfer of electrons between two species.

An increase in oxidation state corresponds to oxidation, while a decrease corresponds to reduction. It's essential to recognize that in redox reactions, one species is oxidized while another is reduced—this is known as the complementary nature of redox processes. A solid grasp of chemical compounds oxidation is important for balancing redox reactions, predicting reaction products, and understanding reaction mechanisms.

For example, when assigning the oxidation state of nitrogen in various compounds, we can discern the context in which nitrogen can gain or lose electrons, leading to different oxidation states. This forms the basis for characterizing the behavior of nitrogen in different chemical environments.

Balancing Oxidation Charges

Balancing oxidation charges involves ensuring that the sum of oxidation states for all atoms in a compound or a molecule is equal to the overall charge of that species. This principle is critical for the conservation of charge and for the stoichiometric balance of chemical equations.

For students to successfully use this concept, familiarizing themselves with the common oxidation states of various elements is necessary, and they must be able to apply the rules for assigning oxidation numbers efficiently. In ionic compounds, balancing charges is relatively straightforward as the charges on the ions should balance out to zero for a neutral compound, or to the net charge for polyatomic ions.

In case of polyatomic ions like ${\mathrm{NO}}_3^{-}$ or ${\mathrm{NO}}_2^{-}$, it's necessary to take into account the known oxidation states of oxygen and any available charges to deduce the oxidation state of the central atom, like nitrogen in these cases. This approach ensures that students can predict and balance the charges consistently across different chemical species.